An atom can be defined as the smallest particle of an element that retains the chemical properties of that element. The particle theory of matter was supported as early as 400 B.C. by Greek thinkers, such as Democritus. He used the term atom based on the Greek word atomos, meaning "indivisible".
Dalton's Atomic Theory
In 1808, John Dalton, an English schoolteacher, provided an explanation for the law of conservation of mass, the law of definite proportions, and the law of multiple proportions (see below in, "Important Terms"). He said that elements were composed of atoms and that only whole numbers of atoms can combine to form compounds
1. All matter is composed of extremely small particles called atoms.
2. Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties.
3. Atoms cannot be subdivided, created, or destroyed.
4. Atoms of different elements combine in simple whole-number ratios to form chemical compounds
5. In chemical reactions, atoms are combined, separated, or rearranged.
Later, however, step 3 of this theory was found to be incorrect.
The Structure of the Atom:
Atoms consist of two regions. One region is the nucleus. The nucleus is a small, dense area located at the center of the atom made up of protons, which are positively charged, and neutrons, which are neutral (no charge), and a region occupied by electrons, which are negatively charged . Protons, neutrons, and electrons are known as subatomic particles.
Investigations of the relationships between matter and electricity with cathode-ray tubes led to the discovery of the electron. Scientists noticed that when current was passed through a cathode-ray tube, the surface of the tube directly opposite the cathode glowed. These tests revealed the following information:
Cathode rays were deflected by a magnetic field in the same manner as a wire carrying electric current, which was known to have a negative charge
The rays were deflected away from a negatively charged object.
These observation led to the hypothesis that the particles that the particles that compose cathode rays are negatively charged. Thus, the electron had been found!
Counting Atoms:
Atomic Numbers:
All atoms are made up of the same basic particles. Even though this is true, not all atoms are the same. As you move from left to right on the periodic table, the atomic number increases.
Isotopes:
Three types of hydrogen atoms are known. They are protium, deuterium, and tritium. These are all isotopes of hydrogen. Isotopes are atoms of the same element that have different masses.
Mass Number:
The mass number is the total number of protons and neutrons that make up the nucleus of an isotope.
Relative Atomic Masses:
The standard used by scientists to compare units of atomic mass is the carbon-12 atom. The mass is exactly 12 atomic mass units, or 12 amu. One atomic mass unit, or 1 amu, is exactly 1/12 the mass of a carbon-12 atom. The atomic mass of any atom is determined by comparing it with the mass of the carbon-12 atom.
Relating Mass to Numbers of Atoms:
The Mole:
The mole is the SI unit for a specific amount of any given substance. A mole is the amount of a substance that contains as many particles as there are atoms in exactly 12 g of carbon-12. The mole is just a counting unit. Do not let it confuse you. The mole is your friend.
Avogadro's Number:
Many different numbers have been chosen to express the number of particles in a mole. The best modern value is 6.022 x 1023 . This means that 12 g of carbon-12 contains 6.022 x 1023 carbon-12 atoms. Avogadro's number is the number of particles in exactly one mole of a pure substance.
Molar Mass:
The mass of one mole of a pure substance is called the molar mass. Molar mass is usually written in units of g/mol. The molar mass of an element is numerically equal to its atomic mass in atomic mass units. For example, O (Oxygen) is 16 g/mol. This can all be found on the periodic table.
Important Terms
Law of conservation of mass: mass is neither created nor destroyed during ordinary chemical reactions or physical changes
Law of definite proportions: a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound
Law of multiple proportions: if two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers
People to know John Dalton: developed the atomic theory
Robert A. Millikan: measured the charge of the electron with the oil-drop experiment
J.J Thomson: discovered the electron with the cathode ray tube experiment.
Ernest Rutherford: found the nucleus of an atom. Also discovered that the atom is mostly empty.
Mole Conversion Practice
(work at the bottom of the page if needed)
1: Convert 12 grams of
NO₂ into moles
2: Convert 2 moles of
H₂O into grams
3: Use the balanced equation:
H₂CO₃ + H₂OH ---->CO₃ + H₃O,
If you had 18 grams of H₂CO₃, how many grams of HCO₃ would be produced after the reaction has occurred?
Funzies:
the path of the electron beam is curved by the magnetic field
Find the molar mass of nitrogen dioxide by adding up the atomic masses of N and O (remember, there are two O's. you should get your molar mass to be about 46 grams)
Now that you've found the molar mass of nitrogen dioxide, divide your 12 grams by the molar mass (meaning, simply divide 12 by 46)
You should get about .26 moles of NO₂
Number 2:
Start with your 2 moles of H₂O
Again, find the molar mass of your element (in this case, H₂O) (you should get 16 grams)
Because your are now finding how many grams are in 2 moles of H₂O, you multiply by the molar mass
You should get about 36 grams of H₂O
Number 3:
Start with your 18 g of H₂CO₃
Find the molar mass of H₂CO₃. You should get 62 g/mol
Divide 18 g by 62 g/mol. You should get about .29 mol
*this is the tricky part* you now have to find the ratio of H₂CO₃ to HCO₃. You can see from the balanced equation that there is only 1 mol of HCO₃ and 1 mol of H₂CO₃. Now you divide .29 mol by 1 mol of H₂CO₃, then multiply by 1 mol of HCO₃. Obviously, you will still have .29
Next, you must find the molar mass of HCO₃. You should get 61 g.
Lastly, multiply .29 by 61 g.
You should get 17.69 g/mol. Meaning, with your 18g of H₂CO₃, you can produce 17.69g of HCO₃.
CHAPTER 3
Atoms: The Building Blocks of Matter
An atom can be defined as the smallest particle of an element that retains the chemical properties of that element. The particle theory of matter was supported as early as 400 B.C. by Greek thinkers, such as Democritus. He used the term atom based on the Greek word atomos, meaning "indivisible".
Dalton's Atomic Theory
In 1808, John Dalton, an English schoolteacher, provided an explanation for the law of conservation of mass, the law of definite proportions, and the law of multiple proportions (see below in, "Important Terms"). He said that elements were composed of atoms and that only whole numbers of atoms can combine to form compounds
1. All matter is composed of extremely small particles called atoms.
2. Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties.
3. Atoms cannot be subdivided, created, or destroyed.
4. Atoms of different elements combine in simple whole-number ratios to form chemical compounds
5. In chemical reactions, atoms are combined, separated, or rearranged.
Later, however, step 3 of this theory was found to be incorrect.
The Structure of the Atom:
Atoms consist of two regions. One region is the nucleus. The nucleus is a small, dense area located at the center of the atom made up of protons, which are positively charged, and neutrons, which are neutral (no charge), and a region occupied by electrons, which are negatively charged . Protons, neutrons, and electrons are known as subatomic particles.
Investigations of the relationships between matter and electricity with cathode-ray tubes led to the discovery of the electron. Scientists noticed that when current was passed through a cathode-ray tube, the surface of the tube directly opposite the cathode glowed. These tests revealed the following information:
These observation led to the hypothesis that the particles that the particles that compose cathode rays are negatively charged. Thus, the electron had been found!
Counting Atoms:
Atomic Numbers:
All atoms are made up of the same basic particles. Even though this is true, not all atoms are the same. As you move from left to right on the periodic table, the atomic number increases.
Isotopes:
Three types of hydrogen atoms are known. They are protium, deuterium, and tritium. These are all isotopes of hydrogen. Isotopes are atoms of the same element that have different masses.
Mass Number:
The mass number is the total number of protons and neutrons that make up the nucleus of an isotope.
Relative Atomic Masses:
The standard used by scientists to compare units of atomic mass is the carbon-12 atom. The mass is exactly 12 atomic mass units, or 12 amu. One atomic mass unit, or 1 amu, is exactly 1/12 the mass of a carbon-12 atom. The atomic mass of any atom is determined by comparing it with the mass of the carbon-12 atom.
Relating Mass to Numbers of Atoms:
The Mole:
The mole is the SI unit for a specific amount of any given substance. A mole is the amount of a substance that contains as many particles as there are atoms in exactly 12 g of carbon-12. The mole is just a counting unit. Do not let it confuse you. The mole is your friend.
Avogadro's Number:
Many different numbers have been chosen to express the number of particles in a mole. The best modern value is 6.022 x 1023 . This means that 12 g of carbon-12 contains 6.022 x 1023 carbon-12 atoms. Avogadro's number is the number of particles in exactly one mole of a pure substance.
Molar Mass:
The mass of one mole of a pure substance is called the molar mass. Molar mass is usually written in units of g/mol. The molar mass of an element is numerically equal to its atomic mass in atomic mass units. For example, O (Oxygen) is 16 g/mol. This can all be found on the periodic table.
Important Terms
People to know
John Dalton: developed the atomic theory
Robert A. Millikan: measured the charge of the electron with the oil-drop experiment
J.J Thomson: discovered the electron with the cathode ray tube experiment.
Ernest Rutherford: found the nucleus of an atom. Also discovered that the atom is mostly empty.
Mole Conversion Practice
(work at the bottom of the page if needed)
1: Convert 12 grams of
NO₂ into moles
2: Convert 2 moles of
H₂O into grams
3: Use the balanced equation:
H₂CO₃ + H₂OH ---->CO₃ + H₃O,
If you had 18 grams of H₂CO₃, how many grams of HCO₃ would be produced after the reaction has occurred?
Funzies:
the path of the electron beam is curved by the magnetic fieldMoles :D
of course, the 10 commolements
mole dictionary and jokes
more jokes!
How To:
Number 1:
- Start with your 12 grams of nitrogen dioxide
- Find the molar mass of nitrogen dioxide by adding up the atomic masses of N and O (remember, there are two O's. you should get your molar mass to be about 46 grams)
- Now that you've found the molar mass of nitrogen dioxide, divide your 12 grams by the molar mass (meaning, simply divide 12 by 46)
You should get about .26 moles of NO₂Number 2:
- Start with your 2 moles of H₂O
- Again, find the molar mass of your element (in this case, H₂O) (you should get 16 grams)
- Because your are now finding how many grams are in 2 moles of H₂O, you multiply by the molar mass
You should get about 36 grams of H₂ONumber 3:
- Start with your 18 g of H₂CO₃
- Find the molar mass of H₂CO₃. You should get 62 g/mol
- Divide 18 g by 62 g/mol. You should get about .29 mol
- *this is the tricky part* you now have to find the ratio of H₂CO₃ to HCO₃. You can see from the balanced equation that there is only 1 mol of HCO₃ and 1 mol of H₂CO₃. Now you divide .29 mol by 1 mol of H₂CO₃, then multiply by 1 mol of HCO₃. Obviously, you will still have .29
- Next, you must find the molar mass of HCO₃. You should get 61 g.
- Lastly, multiply .29 by 61 g.
You should get 17.69 g/mol. Meaning, with your 18g of H₂CO₃, you can produce 17.69g of HCO₃.